Polarity and van der Waals forces
Polarity and van der Waals forces
- Intermolecular forces act between molecules.
- They are much weaker than the bonds inside substances.
- They start with bond polarity.
Practice
Compared with the bonds inside substances, intermolecular forces are:
Intermolecular forces (between molecules) are far weaker than ionic, covalent or metallic bonds within substances.
Polarity and dipoles
- When two atoms of different electronegativity bond, the electrons sit closer to the more electronegative one.
- This gives a dipole: one end slightly negative ($\delta-$), the other slightly positive ($\delta+$).
- If the dipoles don't cancel, the molecule is polar; if they cancel by symmetry (like $\text{CO}_2$), it is non-polar.
Practice
A polar bond has:
Unequal sharing of electrons gives a dipole: δ− at the more electronegative end, δ+ at the other.
Practice
CO₂ is non-polar even though its bonds are polar because:
CO₂ is linear and symmetric, so the two bond dipoles point opposite ways and cancel.
Van der Waals' forces
- Van der Waals' forces is the general name for intermolecular forces. Two types:
- London dispersion (instantaneous dipole–induced dipole): brief dipoles from moving electrons; act between all molecules and get stronger with more electrons.
- permanent dipole–permanent dipole: between molecules that are always polar.
Practice
London dispersion forces get stronger when a molecule has:
More electrons give larger instantaneous dipoles, so London (dispersion) forces are stronger.
You've got it
Key idea
- intermolecular forces are weaker than the bonds inside substances
- a polar bond has $\delta-$ and $\delta+$ ends; a molecule is polar if its dipoles don't cancel ($\text{CO}_2$ cancels → non-polar)
- London dispersion acts between all molecules (stronger with more electrons); permanent dipole–dipole acts between polar molecules